Mole Concept

Mole, which is symbolized n, and has a unit mol is the counting unit for number of atoms, ions or molecules.  One mole is the amount of substance (either in atoms, ions or molecules) as the number of atoms in 12 g of Carbon-12.  Scientists found out that 1 mole of any substance contains 6.02 x 10²³ particles. This number is called Avogadro's Number (N), in honor of the Italian Scientist Amedeo Avogadro (1776-1856) with the unit mol^-1.  Which means there are 6.02 x 10²³ particles per mole of a substance.

Example:

1 mole of C atoms = 6.02 x 10²³ atoms
1 mole of CO₂ molecules =  6.02 x 10²³ molecules
1 mole of SO₃^-2 ions  =  6.02 x 10²³ ions

Mole is just the same as dozen, 1 dozen of eggs = 12 pcs of eggs, 1 dozen of chairs = 12 pcs of chairs.  Anything that is 1 dozen always contains 12 pcs.  One mole on the other hand,  always contains 6.02 x 10²³ particles.

Interconverting  Moles to the number of Particles

Using the Avogadro's number we can determine the number of atoms, ions or molecules of a given substance.

1.  Calculate the number of C atoms in 0.350 mol of C₆H₁₂O₆.  

Solution:

2.  How many moles of CO₂ are in 3.1 x 10²⁴ CO₂ molecules?

Solution:

3.  Aluminum is a metal with a high strength-to-mass ratio  and a high resistance to corrosion.  Calculate the number of atoms of 0.371 mol of Al.

Solution:

4.  Cobalt (Co) is a metal that is added to street to improve its resistance to corrosion.  Calculate the number of moles of 5.00 x 10²⁰ atoms of Co.

Solution:

5.  How many oxygen atoms are in a) 0.25 mol Ca(NO₃)₂, and  b) 1.50 mol of Na₂CO₃?

a.

b.  

Intermolecular Forces of Attraction

Intermolecular forces are attractive forces between molecules.  This forces are responsible  for the non-ideal behavior of gases at high pressures and low temperatures.  This forces are also important in establishing  the form and behavior of matter.

What is the difference between intermolecular forces and intramolecular forces?  Intermolecular forces are forces that holds molecules together while intramolecular forces are hold atoms together in a molecule.  Intramolecular forces are the chemical bond, the covalent, ionic, and metallic.  When intermolecular forces is broken only change in phase occur but when intramolecular forces are broken chemical change occur.  With regards to the strength of these forces, generally intermolecular forces are weaker than intramolecular forces.

The different types of intermolecular forces are hydrogen bond, dipole-dipole, and London  dispersion forces (Dipole-dipole and London dispersion forces are also known as Van der Waals Forces).  This is after Joannes Van der Waals, who developed the equation for predicting the deviation of gases from ideal behavior.  All intermolecular forces are electrostatic which involve the attraction of positive and negative species like ionic bond but weaker than the ionic bond.

Hydogen Bond

Hydrogen bond is a special kind of dipole-dipole interaction.  This exist when hydrogen is attracted to highly electronegative atoms like N, O, F.  Example of this are the polar molecules of NH₃, H₂O and HF.

Above photo shows the hydrogen bond exist in water molecules.  This bond is quite stronger than dipole-dipole interaction due to some reasons.  One is that hydrogen is covalently bonded to high electronegative atoms which lead to special polar X - H bond.  Second is due to small size of hydrogen atom, the dipoles come close together and produce strong dipole-dipole interactions.  

Hydrogen bond are possible only with hydrogen containing compounds because all atoms other than H have inner-shell electrons to shield their nuclei from attraction by lone-pair electrons of the nearby atoms.  Only F, N and O only meet the hydrogen bonding formation due to their high electronegativity value.

Dipole -dipole interaction

Dipole-dipole forces are attractive forces that exist between polar molecules.  Polar molecules are molecules with dipole moments.  The larger the dipole moment the greater the force.  Below shows what happens during dipole-dipole interaction between molecules.  Molecules are arranged in a way that their positive pole is attracted to the negative pole of the other molecule.  

Example of molecules with dipole-dipole interactions are CO, HCl, PCl₃, HBr.

Dispersion forces

Dispersion forces are forces that exist between nonpolar molecules.  Examples of nonpolar molecules are the diatomic molecules, CO₂, CH₄.  This was proposed by Fritz London, a German-American physicist in 1930.  London recognized that the motion of electrons in an atom or molecule can create an instantaneous, or momentary, dipole moment. The strength of dispersion forces tends to increase with increasing atomic or molecular size. Below is an example of dispersion forces between Cl₂ molecules.

Molecular Geometry

Molecular geometry is the three-dimensional arrangement of atoms in a molecule.  This is also called molecular shape.  To determine the molecular geometry of molecule you have to know how to write lewis structure of different molecules.  Molecular geometry of molecule affects the physical and chemical properties such as melting point, boiling point, density and the types of reactions.

In predicting molecular geometry of molecules, VSEPR (Valence-Shell Electron-Pair Repulsion) Model is used.  VSEPR Model accounts for the geometric arrangements of electrons pairs around the central atom in terms of electrostatic attraction between electron pairs. Valence shell means the outermost shell occupied by the valence electrons, responsible for the bonding. 

There are rules to be followed in predicting molecular geometry using VSEPR Model.

1.  Double bonds and triple bonds can be treated as single bond.  However, multiple bonds are larger than single bond, and the electron densities occupies more space.

2.  If a molecule has two or more resonance structures, we can apply VSEPR Model to any one of them.  Formal charges are usually not shown.

In my discussions below, we will be using letters which will represent the central atom,  bonded atoms and lone pairs.  For the central atom, we will use C for easy recall, B for bonded atoms and L for lone pairs.

 Molecules without Lone Pairs on the Central Atom

The molecules that will be mention here are those molecules that do not contain lone pairs or non-bonding electron pairs on the central atom.  And so we will just be using C and B letters, where C stands for the central atom while B the bonded atom.  The pattern to predict the molecular geometry of molecules without lone pairs are CB₂, CB₃, CB₄, CB₅ and CB₆

1.  CB₂ means that there is only one central atom and two bonded atoms, example of this is Berrylium Chloride (BeCl₂).  The lewis structure of this molecule is shown below.

Since there is no lone pair on the central atom, the bonding pairs repel each other, forming a straight line having an angle of 180^o.  The molecule is considered to be linear as shown in the model below:

2.  CB₃ is another pattern which means 1 central atom and 3 bonded atoms without lone pair. Example molecule having this pattern is Boron triflouride (BF₃).  The lewis structure of this molecule is shown below:

The geometry of the molecule having CB₃ pattern is triangular planar, where the angle between bonds is 120^o.  The angle is equal because there is no non-bonding electron pair between the atom.  See the ball and stick model below:

3.  CB₄, is a pattern which shows molecule having 1 central atom and 4 bonded atoms without the lone pair.  Example of a molecule with this pattern is CH₄, in this molecule the carbon atom is bonded to four hydrogen atoms. To show the bonding and non-bonding electrons, lewis structure must be written.  Below is the lewis structure:

Lewis structure shows that there are four bonded atoms and no non-bonding electron pair on the central atom and therefore the angle between bonds will be 109.5^o with a tetrahedral molecular geometry as shown below:

      

      

4.  Another pattern with no non-bonding electrons on the central atom is CB₅.  Example molecule is PCl₅ with the lewis structure below.  

Lewis structure shows that there is only 1 central atom the P (Phosphorus) bonded to 5 Chlorine atoms.  Central atom has no non-bonding electrons and therefore the possible geometry of this kind of molecule is trigonal bipyramidal having 120^o and 90^o between bonds as shown below:

5. CB₆ is another pattern with no non-bonding electron pair or lone pair  on the central atom.  One example of this is SF₆ in which the central atom sulfur is bonded to 6 flourine atoms.  The lewis structure of this molecule is 

From the Lewis structure, there is one central atom the sulfur and 6 surrounding atoms.  Since there is no non-bonding electrons on the central atom the probable molecular geometry of SF₆ is octahedral as shown below:

  

 

SUMMARY OF MOLECULAR GEOMETRY OF MOLECULES WITHOUT LONE PAIR ON THE CENTRAL ATOM

Molecular Geometry of Molecules with the lone pair on the central atom.

The geometry of molecules with lone pair on the central atom is more complicated than those without the lone pair.  These kinds of molecules have three types of repulsive forces that we have to consider:
1. between bonding pairs
2. between lone pairs
3. between bonding pair and lone pair

VSEPR Model predicted the arrangement of repulsive forces in decreasing order:

           lone pair vs lone pair > lone pair vs. bonding pair > bonding pair vs. bonding pair

So here we will use the same symbol for the pattern, C for central atom, B for bonded atoms and L for lone pair.

1.  CB₂L, is the simplest pattern, which means 1 central atom, two bonded atoms with 1 lone pair.  Example of this is SO₂.  The lewis structure is shown below:

The lewis structure shows double bond between oxygen and sulfur, and VSEPR consider this as if single bond.  It has also lone pair on the central atom in which this lone pair exhibits repulsive force with the bonding pair.  Since the lone pair to bonding pair repulsive force is greater than that of  bonding to bonding pair, the angle between O-S-O will be less than 120^o.  And so the molecular shape of this kind of molecule is bent as shown below:

The molecular model of SO₂ is shown below:

2.  CB₃L is another pattern having 1 central atom, three bonded atoms and 1 lone pair.  Example of this kind is NH₃.  Nitrogen is bonded to 3 hydrogen atoms and 1 lone pair. 

The Lewis structure above shows the bonding atoms and the lone pair.  The molecular shape of this kind of geometry is pyramidal, because it looks like a pyramid.  The angle between H-N-H bond will be less than 109.5^o.  Examine the illustration below:

 In the tetrahedral molecular shape the angle is 109.5^o, but once the bonded atom is removed and is changed with a lone pair there is a decrease in the H-N-H bond angle due to the repulsion between the bonding pair and the lone pair.  The angle decreases to 107.3^o.

3.  Another pattern is CB₂L₂, which means molecule having 1 central atom, 2 bonded atoms and 2 lone pairs.  Example of this molecule is the water, H₂O molecule. Look at the lewis structure below:

Oxygen is the central atom being the atom which has smaller number of molecule, surrounded by 2 hydrogen atoms.  There are also 2 lone pairs on the central atom.  The overall arrangement of electron pairs in water is tetrahedral just like NH₃, only that there are two lone pairs.  These two lone pairs tend to be far apart from each other as possible the tendency the O-H bonding pair is push towards each other decreasing the angle to 104.5^o, as shown below:

Therefore the shape of water molecule is bent.

4.  CB₄L is a pattern which means 1 central atom is bonded to 4 atoms with 1 lone pair, like the SF₄ molecule.  Below is the lewis structure of SF₄:

This lewis structure shows that S being the central atom is surrounded with 4 F atoms and 1 lone pair.  The observe shape experimentally is the seesaw,  as shown below:

Another example is the XeF₄, 

5.  CB₃L₂ is another pattern for a molecule having lone pair on the central atom.   Example of this is ClF₃.  The lewis structure is shown below:

As shown above the Cl is the central atom surrounded by 3 bonded atoms the F, and there are 2 lone pairs on the central atom.  The possible molecular geometry of this molecule is T-shaped, as shown below:

6.  CB₂L₃ is another pattern without lone pair on the central atom. This pattern means that there is 1 central atom, 2 bonded atoms and 3 lone pairs on the central atom,  Example molecule is XeF₂.  Let us look at the lewis structure of this molecule. 

The possible shape of this molecule is  linear, as shown in the model below:

7.  CB₅L, a pattern for a molecule having 1 central atom, 5 bonded atoms with 1 lone pair in the central atom.  Example of this is BrF₅,  

Lewis structure above shows that Br is surrounded by 5 F atoms with 1 lone pair.  The molecular shape of this kind of molecule is  square pyramidal as shown below:

8.  CB₄L₂, is another pattern showing molecule with 1 central atom, 4 bonded atoms with 2 lone pairs,  Example molecule is XeF₄ with the lewis structure of 

Based from the VSEPR the possible shape of this kind of molecule is square planar, as shown below:

These are the different molecular geometries of molecule having lone pair in the central atom.  

RELATIONSHIP BETWEEN MOLECULAR GEOMETRIES

1.  With 2- 4 number of electron domains

2.  With 5 number of electron domains

3.  With 6 number of electron domains

Bond Polarity

Bond Polarity is used to describe the sharing of electrons between atoms in covalent bonding.  A polar covalent bond is a bond in which there is unequal sharing of electrons between atoms while a nonpolar covalent bond has equal sharing of electrons between atoms.

How can we distinguish between polar and nonpolar bond?

Electronegativity is a property of the elements used to determine if the sharing of electrons between atoms in covalent bonding is equal or not equal.  Electronegativity is defined as the ability of an atom in a molecule to attract electrons to itself.  The greater the electronegativity value the greater its ability to attracts electrons to itself.

The difference in electronegativity between atoms in a molecule is used to predict the sharing of electrons.  Consider the example below:

Example 1.

Compound   :         F₂

Electronegativity difference :     4.0 - 4.0  =  0

Type of Bond :    Nonpolar covalent bond



Example 2

Compound   :  HF

Electronegativity difference :  4.0 -  2.1 =  1.9

Type of Bond    :   Polar covalent bond



Example 3.

Compound   :    LiF

Electronegativity Difference:  4.0 - 1.0  = 3.0

Type of Bond  :    Ionic Bond



The three examples above show how the type of bond be identified.  In example 1,  the difference between the electronegativity of two F atoms is 0, which is an indication that the sharing of electrons between atoms are equal.  Therefore, a nonpolar covalent bond results when the electronegativity difference between atoms is from 0 to 0.4.  In this difference of electronegativity still shows that the sharing of electrons is equal between atoms.

In example 2, we can see that the difference of electronegativity values between atoms is 1.9, this shows that the sharing of electrons is not equal.  Therefore, a polar covalent bond results when the electronegativity difference between atoms ranges from 0.5 to 1.9.   At this difference the sharing of electrons between atoms resulted in unequal sharing.

In example 3, the difference of electronegativity values between atoms is 3.0 which means that there is a big difference of electronegativity which resulted the transfer of electrons between one atom to another atom.  Therefore, Ionic bond exist when the electronegativity difference is from 2.0 above.

Covalent Bond

Covalent bond is another type of bond that involves the sharing of valence electrons between nonmetal atoms forming a molecule.  Covalent compounds are the compounds that contains only covalent bonds. Examples of this are the diatomic molecules:  H₂, O₂, N₂.

To show the sharing of electrons between atoms in a molecule lewis structure must be written.  A Lewis structure is a representation of covalent bonding in which shared electron pairs are shown either as line or pair of dots, and lone pairs are shown as pair of dots on the individual atoms.  A line is used when electron pair is shared between atoms and dots are the unshared pair electrons.

In writing the Lewis structure, octet rule must also be followed and with a few exemptions.

Lewis Structure

There are rules to be followed when writing Lewis structure:
1.  Add the total valence electrons from all atoms.  For polyatomic anions, add the negative charges to the total number of valence electrons and for the polyatomic cations, subtract the positive charges from the total number of valence electrons.

2.  Write the symbols of all the elements that comprise the molecule, arranging them in such a way that the central atom is the least electronegative atom or the atom with the least number of atoms.  Example for the molecule CF₄, C is the central atom having only one atom and F  are the bonded atoms having 4 F atoms, F also is the most electronegative atom.

3.  Draw a single covalent bond between the central atom and the surrounding atoms.  Complete the 8 valence electrons of the bonded atoms and the excess valence electrons will be placed to the central atom and will represent the non-bonding electrons.

4.  After completing step 1 - 3, if the central atom did not complete 8 number of valence electrons, try adding double bond or triple bond between the surrounding atoms and the central atoms.


Examples 1

Write the Lewis structure of the NF₃
Step 1.  Add the total number of valence electrons.  Nitrogen is located in group VA therefore has 5 valence electrons and Flourine is in group VIIA , 7 valence electrons.

Total valence electrons:   5 + 3(7) = 5  +  21  = 26 valence electrons

Step 2.  Arranging the atoms. the least electronegative atom is the central atom (N)  and  F are the bonded atoms.

Step 3. Draw a single covalent bond between the central atom and the surrounding atoms. Complete the 8 valence electrons of the bonded atoms and the excess valence electrons will be placed to the central atom and will represent the non-bonding electrons.

Step 4.  No more step 4 since all the atoms in the compound achieved stability by having eight valence electrons.

Example 2. 

 Write the Lewis structure of the carbonate ion, CO₃^-2

Step 1.  Calculate the total valence electrons.  C has 4 valence electrons (IVA) and O has 6 valence electrons (VIA).

Total valence electrons:   4 +  3 (6)  +  2 = 4  +  18  +  2  =  24 valence electrons
                                       

Step 2.  Arranging the atoms

Step 3.  Draw a single covalent bond between the central atom and the surrounding atoms. Complete the 8 valence electrons of the bonded atoms and the excess valence electrons will be placed to the central atom and will represent the non-bonding electrons.

Step 4.  Since the central atom carbon is not yet stable, the two electrons from either of the oxygen can be made into double bond so the eight valence electrons in carbon will be meet. As shown below:

Example 3.

Draw the lewis structure of  F₂ gas.

Step 1.  Calculate the total valence electrons of F_2.   2(7) =  14 valence electrons

Step 2.  Arranging the atoms.  Since there are only 2 electrons, you can arrange them in linear form.

Step 3.  Distribution of valence electrons, first a single bond will be placed in between the two fluorine atoms, the rest of electrons will be placed surrounding the F atom.

Step 4.  Since each F atom, already achieved 8 valence electrons the structure above is already the lewis structure.  If in case each atom will not achieve 8 valence electrons you can make double bond or triple bond between atoms.

TRY THIS:

Write the Lewis structure of the following:

1.  SF₆

2.  PO₃^-3

3. CCl₄

4.  CO₂

5.  H₂S